Redox potential

The essence of conjugated redox pairs

At the heart of any redox reaction is the interaction between two conjugated systems. The process reduces to the mutual conversion of substances: the oxidizing agent, by accepting electrons (Reduction), transforms into the conjugated reducing agent—the reduction product. At the same time, the original reducing agent, by releasing electrons (Oxidation), converts into the conjugated oxidizing agent form.

For example, during the interaction of iron(III) and tin(II) ions, iron acts as an oxidizing agent and converts into a divalent cation. Tin, conversely, donates electrons and increases its oxidation state, becoming a conjugated oxidizing agent. As a result, two independent pairs are formed that exchange charges.

How standard potential works

The quantitative measure of each pair's activity is the standard redox potential ($E^\circ$). Its physical meaning lies in assessing how actively a substance attracts electrons.

The pattern is simple:

• The higher the $E^\circ$ value, the more pronounced the oxidizing properties (and, accordingly, the weaker the reducing agent formed from it).
• The lower the $E^\circ$ value, the more active the reducing agent (and the weaker its conjugated oxidizing agent).

For example, the standard potential of the iron pair is $+0.77$ V, and that of the tin pair is $+0.15$ V. Since the value for iron is higher, it is iron that accepts electrons from tin, acting as the dominant oxidizing agent.

Rule for the direction of chemical reactions

Knowing the potentials allows predicting whether a reaction will occur. The main rule for spontaneity states that a strong oxidizing agent and a strong reducing agent react with each other to form weak conjugated products. Electrons always move from the system with a lower potential value to the system with a higher one.

As the charge transfer proceeds, the concentration of the initial active substances decreases, their action weakens, and the system gradually reaches chemical equilibrium. Meanwhile, the completeness of the reaction directly depends on the difference in standard potentials ($\Delta E^\circ$) of the two interacting pairs: the greater this difference, the more thorough the conversion of the initial reactants into products.

Selecting reactants using tables

When planning laboratory or industrial syntheses, the necessary substances are selected based on tables of standard electrode potentials.

• For oxidation, reagents from the upper part of the tables are chosen. Typical powerful oxidizing agents include halogens, hydrogen peroxide, ozone, permanganates, and dichromates. For example, to oxidize hydrogen sulfide to free sulfur, tin salts are not suitable (they are located lower than sulfur in the table). Oxygen has the required potential, but reacts extremely slowly, making molecular iodine the optimal choice.
• For reduction, one refers to the lower part of the list. Active metals, tin(II) salts, sulfur dioxide (sulfur(IV) oxide), and sulfites are located there. Thus, to obtain a bromide from a bromate, using potassium sulfite is an excellent option, as it reliably reduces the halogen to the required oxidation state.

Class and exam questions

What else to study on this topic

• The Nernst equation and the effect of concentration on potential
• The relationship between redox potential and Gibbs free energy
• Kinetic limitations of reactions (using dissolved oxygen as an example)
• Galvanic cells and calculating electromotive force (EMF)

Other topics in "Chapter 7. Classification of chemical reactions by electronic processes"

Prepared by WhiteCoat Pro faculty in accordance with the medical school curriculum. For educational purposes only.